A sample of 0.9535 g of pure borax, Na₂B₄O₇· 10H₂O (molar mass 381.37 g/mol), is dissolved in water…

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A sample of $0.9535\ \text{g}$ of pure borax, $\text{Na}_{2}\text{B}_{4}\text{O}_{7}\cdot 10\text{H}_{2}\text{O}$ (molar mass $381.37\ \text{g/mol}$ ), is dissolved in water and titrated with hydrochloric acid of concentration $0.0500\ \text{mol/L}$ . The reaction is $\text{Na}_{2}\text{B}_{4}\text{O}_{7} + 2\text{HCl} + 5\text{H}_{2}\text{O} \to 2\text{NaCl} + 4\text{H}_{3}\text{BO}_{3}$ . Find the volume of acid needed to reach the end point. Then find the mass of boric acid ( $61.83\ \text{g/mol}$ ) and of sodium chloride ( $58.44\ \text{g/mol}$ ) that form.

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Borax reacts with acid as $\text{Na}_{2}\text{B}_{4}\text{O}_{7} + 2\text{HCl} + 5\text{H}_{2}\text{O} \to 2\text{NaCl} + 4\text{H}_{3}\text{BO}_{3}$ . Heating gives $\text{Na}_{2}\text{B}_{4}\text{O}_{7}\cdot 10\text{H}_{2}\text{O} \to \text{Na}_{2}\text{B}_{4}\text{O}_{7} \to 2\text{NaBO}_{2} + \text{B}_{2}\text{O}_{3}$ . For a weak acid, $[\text{H}^+] = \sqrt{K_a c}$ , and a mixture of the acid and its conjugate base obeys $\text{pH} = \text{p}K_a + \log_{10}\dfrac{[\text{base}]}{[\text{acid}]}$ . By Hess's law, $\Delta H = \sum \Delta_f H(\text{products}) - \sum \Delta_f H(\text{reactants})$ .